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Chapter 6 Lecture 1 Acid-Base Concepts

Chapter 6 Lecture 1 Acid-Base Concepts. Unifying Concepts The Acid-Base Concept There are many acid-base definitions, each at times useful Acid-Base concepts are not facts or even theories, but are useful generalizations for classification, and organization

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Chapter 6 Lecture 1 Acid-Base Concepts

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  1. Chapter 6 Lecture 1 Acid-Base Concepts • Unifying Concepts • The Acid-Base Concept • There are many acid-base definitions, each at times useful • Acid-Base concepts are not facts or even theories, but are useful generalizations for classification, and organization • Acid-Base concepts are powerful ways to explain data and predict trends • Arrhenius Concept • An acid forms H+ in water; a base forms OH- in water • Applicable to aqueous solutions only • HCl + NaOH H+ + OH- + Na+ + Cl- • Bronsted-Lowery Concept • Acid is a proton donor; Base is a proton acceptor • Conjugate acid/base pairs differ only by a proton • Reactions proceed to produce the weakest acid and base

  2. H3O+ + NO2- H2O + HNO2 • Includes non-aqueous systems NH4+ + NH2- 2 NH3 • Solvent System Concept • Useful for aprotic, non-aqueous systems • Applies to any solvent that can dissociate to cation (acid) and anion (base) • For water: 2 H2O H3O+ + OH- • Any solute increasing [H3O+] is an acid HCl + H2O H3O+ + Cl- • Any solute increasing [OH-] is a base NH3 + H2O NH4+ + OH- • Aprotic, non-aqueous example: 2 BrF3 BrF2+ + BrF4- • Acid: SbF5 + BrF3 BrF2+ + SbF6- • Base: F- + BrF3 BrF4- • The acid-base reaction: acid + base = solvent (reverse of the ionization) • Arrhenius: acid + base = salt + water • Bronsted: acid1 + base2 = base1 + acid2

  3. pKion = -log[acid][base] • pKW = -log[H3O+][OH-] = -log[10-7][10-7] = 14 • pKH2SO4 = -log[H3SO4+][HSO4-] = 3.4 • The smaller the number, the more dissociation has occurred • Lewis Concept • Acid = e- pair acceptor; Base = e- pair donor • Includes metal ions and non-aqueous systems; encompasses other concepts • We will use this concept throughout the rest of the chapter and course • Metal ion Example: Ag+ + 2 NH3 [Ag(NH3)2]+ • Acid-Base product is called an adduct • If the acid is a metal ion, it is also called a coordination compound or a coordination complex or a complex ion

  4. 5) A non-metal example: BF3 + NH3 H3N:BF3 (or BF3• NH3)

  5. Acid-Base Strength • Thermodynamic Measurement • We can easily measure pH, but that doesn’t really tell us about acid strength • HA + H2O H3O+ + A- • DGo = -RTlnKa = DH - TDS • G = free energy • H = enthalpy • S = entropy • Solving for Ka: Ka

  6. Binary Hydrogen Compounds • Acidity increases down a column of the periodic table • H2Se > H2S > H2O • HI > HBr > HCl > HF • Conjugate bases of larger ions have lower charge density, thus a smaller attraction for H+ • Acidity increases from left to right of the periodic table • NH3 < H2O < HF • The more electronegative the conjugate base is, the easier is it for H+ to dissociate

  7. Inductive Effects (electron pulling/pushing through sigma bonds) • Electronegative substituents increase acidity and decrease basicity Basicity: :PF3 < :PH3 • Electron Donating substituents decrease acidity and increase basicity Basicity: NMe3 > NHMe2 > NH2Me > NH3 • Oxyacids: the more unprotonated Oxygens, the stronger the acid • Acidity: HOClO3 > HOClO2 > HOClO > HOCl • The electronegative O’s pull e- away from the H—O bond • The electronegative O’s stabilize the conjugate base • Cations in Aqueous Solution • Cationic metal ions are generally Lewis acids in water solutions • Example: [Fe(H2O)6]3+ + H2O [Fe(H2O)5(OH)]2+ + H3O+ • Large charge and small radii increase acidity • Alkali metals are not acidic (Na+); Alkaline Earths are weakly acidic (Ca2+) • 2+ Transition Metals are weak acids; 3+ Transition Metals are strong acids • All 4+ or higher metals are very strong acids  MxOy

  8. 4) The stronger acid the cation is, the less soluble the hydroxide complex is. OH- can’t dissociate to dissolve because of strong charge attraction. We can use this property to estimate the acid strength of the cation

  9. Steric Effects • Steric bulk can repel an acid-base partner, modifying the acid-base strength • F = front strain = direct steric interference at the site of interaction • B = back strain = bulky groups interfere opposite the interaction site upon binding as the molecule adjusts its VSEPR geometry • The order of basicity can scramble depending on bulk of the acid 4 4 4 3 2 2 2 1 1 1 3 3

  10. Solvation • Solvation is interaction with solvent molecules • Basicity in water: NHMe2 > NH2Me > NMe3 > NH3 • By induction, the more substituted amine should be the most basic • This amine has less H’s to interact with water • Non-aqueous Solvents • The Leveling Effect: the strongest acid possible in a solvent is the solvent cation; the strongest base possible in a solvent is the solvent anion. • H2SO4 + H2O H3O+ + HSO4- (100% dissociation) • Na2O + H2O 2 Na+ + 2 OH- (100% dissociation) • H2SO4 + HOAc H2OAc+ + HSO4- (< 100%) • NH3 + HOAc NH4+ + OAc- (100%) • HNO3, H2SO4, HClO4, HCl are all equally acidic in water (H3O+) • HClO4 > HCl > H2SO4 > HNO3 in HOAc • Hydrocarbon Solvents don’t level acids or bases

  11. Superacids = acids stronger than H2SO4 • Hammet Acidity Function = Ho B = nitroaniline indicator used as the base • Lewis Superacids are often made by protonating an already strong acid • This is often done using HF as the acid to be protonated • It requires a very stable anion to make the reaction proceed • 2 HF + 2 SbF5 H2F+ + Sb2F11- (Fluoroantimonic acid)

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